c- he-m- i= s-tr y for conservators · lead pb ancient latin, plumbum, meaning heavy french,...

1 20-Jul-17 © Prof. Zvi C. Koren

Introduction to 6C-2He-m-53I=16S-tr-39Y

for Conservators

Shenkar College: Engineering. Design. Art.

The Edelstein

Center

www.edelsteincenter.wordpress.com www.shenkar.ac.il

http://www.edelsteincenter.wordpress.com/

http://www.shenkar.ac.il/


1. The Periodic Table of the Elements

Mendeleev

Element names

Natural vs. synthetic elements

Metals, nonmetals, semi-metals (metalloids)

Solids, liquids, gases

Groups (families, columns) and periods (rows)

Molecular elements

Allotropes

Periodicity examples: The alkali metals

2. Atoms

Atomic structure

Atomic number and mass number

Isotopes

Atomic weight/mass

3. Ionic & Covalent Compounds: Formulas

Nature of chemical bonds

Cations & anions

Nomenclature

Acids & bases - names

Empirical vs. molecular formulas

Syllabus 4. Chemical Reactions

Acids and bases - reactions

Reaction types

Balancing reactions

Net ionic reactions

5. Measurement Units & Conversions

Scientific notation

Metric and S.I. system

Unit prefixes

Calculations via the “unit conversion factor” method

6. The Mole

Avogadro’s number

Atomic and molecular weights/masses

Gram-mole-items calculations

7. Stoichiometry: Chemical Mathematics

Chemical quantities reacting & produced

Solution concentration units & conversions:

Molarity (M), molality (m), Normality (N)

Weight percent (w/w%, w/v%); Volume percent (v/v%)

Parts (ppt, ppm, ppb)


1.

The Chemical Bible:

The Periodic Table of the Elements


Dmitri Ivanovich Mendeleev Дмитрий Иванович Менделеев

(1834 –1907)

On 6 March 1869, Mendeleev made a formal presentation

to the Russian Chemical Society, titled:

On the Relationship of the Properties of the Elements to their Atomic Weights. (Actually, Mendeleev was ill, and his colleague Nikolai Menshutkin presented his paper.)

The paper was published in the first volume of the new society's journal:

Journal of the Russian Chemical Society 1, 60-77 (1869):

That same year, a German abstract of the

paper, consisting of the table and eight

comments, was published:

Zeitschrift für Chemie 12, 405-406 (1869).

The German abstract was the vehicle by

which Mendeleev's ideas reached chemists

working in Western Europe.

7 20-Jul-17 © Prof. Zvi C. Koren pp. 405 – 406


Artist: Ivan Nikolaevich Kramskoi (1878)

D.Mendeleev Museum & Archives, St.Petersburg State University

http://spbu.ru/files/culture/museums/mendeleev/index.html

https://upload.wikimedia.org/wikipedia/commons/d/d4/Kramskoy_Mendeleev_01.jpg


Derivation of Name or

Symbol

Discoverer

(Country)

Date of

Discovery

Symbol Element

Berkeley, California (site of Seaborg’s

laboratory)

G.T. Seaborg,

S.G. Thompson,

A. Ghiorso (USA)

1950 Bk Berkelium

Latin, cuprum, derived from Cyprium,

the island of Cyprus, the main source

of copper in the ancient world.

Ancient Cu Copper

Albert Einstein A. Ghiorso (USA) 1952 Es Einsteinium

Latin, ferrum Ancient Fe Iron

Latin, plumbum, meaning heavy Ancient Pb Lead

French, oxygene, generator of acid,

derived from the Greek, oxy and

genes, meaning acid-forming; oxygen

was thought to be part of all acids

J. Priestley (UK),

K.W. Scheele

(Sweden)

1774 O Oxygen

Latin, argentum Ancient Ag Silver

Latin, stannum Ancient Sn Tin

Name from Swedish, tung sten,

meaning heavy stone; symbol from

wolframite, a mineral

J.J. and F. de

Elhuyer (Spain)

1783 W Tungsten

Derivation of Element Names and Symbols – People, Places, and Things


Modern Periodic Table

is based on

“Atomic Numbers”

The Periodic Table Matrix

Periods, Rows

Groups, Columns, Families

Main group elements

Transition group elements

Metals

Nonmetals

Metalloids, Semimetals

Alkali metals

Alkaline earth metals

Coinage metals

Halogens

Noble gases

Lanthanides & Actinides


He H

Ne F O N B

Ar Cl Al

Kr Br Ga

Xe

Rn Hg Cs

Fr

@ room temperature (250C)

@ Room Temperature (all others are solids)

Liquid Gaseous & Elements


H

F O N

Cl

Br

I

At

Stable Elements as Diatomic Molecules (@ Room Temperature)

H2, N2, O2, F2, Cl2

Br2

I2

At2

הנוף

כולו

בר

י

א


The Halogens

diatomic molecules

Molecular Elements


Carbon

Allotropes


Oxygen

Allotropes

triatomic

molecule O2

light 2 O

O + O2 O3

diatomic

molecule


tetrahedral

(tetrahedron)

tetratomic

molecule


octatomic

molecule


Examples of Periodicity:

Reactions of Alkali Metals with Water

https://www.youtube.com/watch?v=uixxJtJPVXk

https://www.youtube.com/watch?v=jI__JY7pqOM

Summary: https://www.youtube.com/watch?v=tubhteKh_dQ

https://www.youtube.com/watch?v=QQF61CFOySw

Li https://www.youtube.com/watch?v=Vxqe_ZOwsHs

Na https://www.youtube.com/watch?v=dmcfsEEogxs

K https://www.youtube.com/watch?v=oqMN3y8k9So

https://www.youtube.com/watch?v=eaChisV5uR0

1A

H

Li

Na

K

Rb

Cs

Fr

Li

Na

K

https://www.youtube.com/watch?v=uixxJtJPVXk




https://www.youtube.com/watch?v=tubhteKh_dQ




https://www.youtube.com/watch?v=Vxqe_ZOwsHs

https://www.youtube.com/watch?v=dmcfsEEogxs

https://www.youtube.com/watch?v=oqMN3y8k9So












1A

H

Li

Na

K

Rb

Cs

Fr

Chemical Periodicity

Li

Na

K

M(s) + H2O(l) MOH(aq) + ½H2(g) + Heat

Reaction Names:

Hydrolysis

Single Replacement

“Redox”

(Reduction-oxidation)

Exothermic

Gas-Forming

aqueous

Elements of the same group have similar – but not identical – properties.


2.

Atoms


Democritus (~ 460 – 370 BCE)

ἄτομος , ἀτόμους , ἄτομα , … atomos = a + tomos = uncut, undivided, indivisible

All matter was eventually reducible to discrete, small

particles or atomos.

“Nothing exists except atoms and

empty space, everything else is

opinion.”

Democritus is known as the

"Laughing Philosopher" because of

his joyous spirit. He was a big man

(relatively speaking) and enjoyed life

tremendously.


amu = atomic mass unit

Atomic Composition

e–

H+

Relative Atomic Mass Scale:

Mass of 1 atom of 12C 12 (exactly) amu


Atomic Symbols & Isotopes (איזוטופים)

XA

ZElement symbol

Atomic Number # of protons =

Mass Number # of protons + # of neutrons =

Examples of Isotopes:

C14

6 C

13

6 C

12

6 :carbonelement theof isotopes Three

for a neutral atom: # of electrons = # of protons

protium deuterium (D) tritium (T)

Isotopes = different atoms of the same element Radioactive

Consider ions (charged species – “to go”): 13C4-: #p = __, #n = __, #e = __. 3H+: #p = __, #n = __, #e = __.

carbon-12 carbon-13 carbon-14 #p =

#n =

#e =

#p =

#n =

#e =

1H 2H 3H Three isotopes of the element hydrogen:


Atomic Mass or Atomic Weight (A.W.)

the weighted average of all the stable isotopes of that element

Atomic Mass or

Atomic Weight (amu)

Percent (%)

Abundance Mass (amu)

Mass

Number Symbol Element

1.00797 99.9855 1.007825 1

H Hydrogen 0.0145 2.014102 2

10.811 19.91 10.012939 10

B Boron 80.09 11.009305 11

15.9994

99.759 15.994914 16

O Oxygen 0.0374 16.999134 17

0.2039 17.999160 18

12.011

12

(exactly) 12

C Carbon

13

Table of Exact Masses of the Stable Isotopes of Some Elements

For example:

A.W. of O = 0.99759(15.994914) + 0.000374(16.999134) + 0.002039(17.999160)

= 15.9994 amu [This is the mass that appears in the Periodic Table.]

?


Notes about Atomic Weights

• All atomic masses (except for 12C) are NOT integers

• Mass number is always an integer!

• A.W. the sum of all the particles.

For example for deuterium, 2H (or D):

1 proton = 1.007276 amu

1 neutron = 1.008665

1 electron = 0.0005485799

Total mass of particles = 2.0164895799 amu

Atomic Weight = 2.014102 amu

Difference in mass = 0.002387 amu = 3.9638x10-30 kg

Why?

Mass defect: E = mc2 = (3.9638x10-30 kg)(2.997925x108 m/s)2 = 3.5625x10-13 J/atom

1 atom = 2.014102 amu = 3.3445x10-24 g

energy/gram = 1.0652x1011 J/g = 1.0652x108 kJ/g

energy/mole = 2.1454x108 kJ/mole

For exothermic chemical reactions: energy/mole 102 kJ/mole


3.

Ionic & Covalent Compounds:

Molecular Formulas


The Language of Chemistry

Letters = Element Symbols

↓ ↓

Words = Compound Formulas


Compounds

Covalent (between Nonmetals)

Ionic (typically between Metals & Nonmetals)

composed of ions

positive (+):

Cation

negative (-):

Anion

Electric

“Glue”

composed of e-sharing

Atoms share the same e-pair

A : B

Salts & Oxides


Ionic Compounds

composed of ions: positive(+) and negative(-)

when Metals meet Nonmetals, they instantly fall in LOVE

Metals LOVE to transfer electrons to Nonmetals

“opposites attract” or “viva la difference”


Octet Rule: Atoms combine to mimic the very stable noble elements

Main Group Number is the number of valence electrons = Number of e’s in outermost shell

Typical charges for monatomic ions:

Na+ Mg2+ Al3+ C4- N3- O2- F-

Examples of ions:

1+ 2+ 3+ 4- 3- 2- 1-

Na + Cl Na Cl + – He

Ne

Ar

Kr

Xe

Rn

X

Ionic Charges

Ca & Br:

Examples of compounds:

Ca2+ + 2Br– CaBr2(s)

Ba & S: Ba2+ + S2– BaS(s)

Al & O: K & N:

2Al3+ + 3O2– Al2O3(s) 3K+ + N3– K3N(s)


Transition Metals & Others (continued)

Fe Cu

Hg

Sn

Pb

Common Ions

with two possible charges

-ous and –ic

Names

Roman-Numeral

Name

Ion

Ferrous ion Iron(II) ion Fe2+

Ferric ion Iron(III) ion Fe3+

Cuprous ion Copper(I) ion Cu+

Cupric ion Copper(II) ion Cu2+

Mercurous ion Mercury(I) ion Hg22+

Mercuric ion Mercury(II) ion Hg2+

Stannous ion Tin(II) ion Sn2+

Stannic ion Tin(IV) ion Sn4+

Plumbous ion Lead(II) ion Pb2+

Plumbic ion Lead(IV) ion Pb4+


Nomenclature of Ionic Compounds (שיטת כינוי, כיניון, מינוח)

Cation: Metal = Element name or

NH4+ = ammonium ion (and its derivatives: CH3NH3

+, …) – polyatomic cation

Anion: Monatomic = Element root name + ide

7A 6A 5A 4A

H– hydride

F– fluoride O2– oxide N3– nitride C4– carbide

Cl– chloride S2– sulfide P3– phosphide

Br– bromide Se2– selenide

I– iodide Te2– telluride

Examples:

CaBr2

BaS

Al2O3

K3N

FeCl2

FeCl3

Hg2Cl2

HgCl2

LiH

(NH4)3P

calcium bromide

barium sulfide

aluminum oxide

potassium nitride

iron(II) chloride = ferrous chloride

iron(III) chloride = ferric chloride

mercury(I) chloride = mercurous chloride

mercury(II) chloride = mercuric chloride

lithium hydride

ammonium phosphide


NOT composed of ions

(but consist of partial charges: δ+ and δ–)

Covalent Compounds (between Nonmetals)


C2H6O

Formulas - Representations

Molecular Formula

Condensed Structural Formula

Expanded Structural Formula

Stereo Projection Formula

Molecular Model

CH3CH2OH

OH

H

H

H

H H

CC

Example: Ethanol or Ethyl Alcohol

HO

H

H

H

H

H

CC

(ball & stick) (spacefill)


Nomenclature of Binary Nonmetal Compounds: AaBb

prefix #

mono 1

di 2

tri 3

tetra 4

penta 5

hexa 6

hepta 7

octa 8

nona 9

deca 10

EXAMPLES:

NF3

NO

NO2

N2O

N2O4

PCl5

SF6

S2F10

IF7

HCl

H2S

H3As

What do these molecules look like?

General Rule for AaBb:

With no H’s: (prefix mono)(Element name for A) (prefix)(name of B as an ide)

With H’s: same as above without any prefixes

H2O water

NH3 ammonia

N2H4 hydrazine

PH3 phosphine

NO nitric oxide

N2O nitrous oxide

Compounds

with

Historical Names

(do not follow the rules):

nitrogen trifluoride

nitrogen monoxide

nitrogen dioxide

dinitrogen monoxide

dinitrogen tetraoxide

phosphorus pentachloride

sulfur hexafluoride

disulfur decafluoride

iodine heptafluoride

hydrogen chloride

hydrogen sulfide

hydrogen arsenide


Substance with an ionizable proton in an aqueous solution: HCl(aq) H+ + Cl–

Naming Acids

• If anion’s suffix is ide, the acid’s name is: hydro(anion root name)ic acid:

Compound(aq), Acid Pure Compound Anion

HCl(aq), hydrochloric acid HCl(g), hydrogen chloride Cl– chloride

HBr(aq), hydrobromic acid HBr(g), hydrogen bromide Br– bromide

H2S(aq), hydrosulfuric acid H2S(g), hydrogen sulfide S2– sulfide

HCN(aq), hydrocyanic acid HCN, hydrogen cyanide CN– cyanide

• For oxyacids (oxoacids): if anion’s suffix is ate, the acid’s name is: (anion root name)ic acid if anion’s suffix is ite, the acid’s name is: (anion root name)ous acid

HNO3(aq) nitric acid NO3– nitrate

HNO2(aq) nitrous acid NO2– nitrite

H2SO3(aq) sulfurous acid SO32– sulfite

H2SO4(aq) sulfuric acid SO42– sulfate

HOCl(aq) hypochlorous acid ClO– hypochlorite

HClO2(aq) chlorous acid ClO2– chlorite

HClO3(aq) chloric acid ClO3– chlorate

HClO4(aq) perchloric acid ClO4– perchlorate

H3PO4(aq) phosphoric acid PO43– phosphate

H3PO3(aq) phosphorous acid PO33– phosphite


Hydrated Compounds


Nomenclature of Selected Compounds

Formula Name

MgCl2 magnesium chloride

Fe(ClO3)3 iron(III) chlorate ALSO ferric chlorate

Fe(ClO2)2 iron(II) chlorite ALSO ferrous chlorite

Mg(SCN)2 magnesium thiocyanate

N2O5 dinitrogen pentoxide

HCl(g) hydrogen chloride

HCl(aq) hydrochloric acid

HClO3(aq) chloric acid

HClO2(aq) chlorous acid


4.

Chemical Reactions


Acids & Bases – Preview

Acid (simple definition) = Substance that produces H+ ions

in water

HF(aq) H+ + F–

Strong Acid (Strong Electrolyte):

Weak Acid (Weak Electrolyte):

Common Strong Acids

HCl < HBr < HI

HClO3

HClO4

HNO3

H2SO4 (1st proton)

Base (simple definition) = Substance that produces OH–

ions in water

Common Strong Bases

MOH, M=Metal from Group I

M(OH)2, M=Metal from Group II

[But solubility limits the basicity]

ammonium ion

hydrochloric acid

hydrofluoric acid

Strong Base (Strong Electrolye):

NaOH(aq) Na+ + OH- sodium hydroxide

Ca(OH)2(aq) Ca2+ + 2OH-

calcium hydroxide

NH3(aq) + H2O(l) NH4+ + OH-

Weak Base (Weak Electrolyte):

ammonia

Moderately Strong Base (Strong Electro.):

acid + base salt + water


+

Acid Base

Salt

Water

The Birth of a Salt: A Family Portrait

Koren’s Acid-Base Family Rule: The Strength of the Baby-Salt

will be according to the Stronger Parent


Acid-Base Properties of Salts – Qualitative Hydrolysis of Salts

Reminder: The rxn between an acid and a base produces a SALT. For example, HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) (Daddy) (Mommy) (Baby)

Types of Salts: (Which are acidic, basic, or neutral?) 1. Salt from a strong acid & a strong base: NaCl 2. Salt from a weak acid & a strong base: NaAc, Na2CO3

3. Salt from a strong acid & a weak base: NH4Cl 4. Salt from a weak acid & a weak base: NH4Ac

Koren’s Family Rule: The nature of the baby is influenced by the stronger parent!


H2O H+ + OH-

Equilibrium constant: Water ionization constant Kw = [H+][OH-] = 1.0 x 10–14 @ 25oC

In neutral solutions and pure water: [H+] = [OH-] = 1.0 x 10–7 M In acidic solutions: [H+] > [OH-] In basic solutions: [H+] < [OH-] but [H+][OH-] = 1.0 x 10–14

The Acid-Base Properties of Water:

Auto-Ionization


Sorensen: “The potential (or power) of H”

pH –log[H+]

(Note: small “p” big “H”)

The pH Scale

@ 25 oC: For neutral solutions: pH = 7.00 For acidic solutions: pH < 7.00 For basic solutions: pH > 7.00

From before: [H+][OH-] = 1.0 x 10–14 @ 25 oC: In neutral solutions and pure water: [H+] = [OH-] = 1.0 x 10–7 M In acidic solutions: [H+] > [OH-] (e.g., [H+] = 10–6, [OH–] = 10–8) In basic solutions: [H+] < [OH-] (e.g., [H+] = 10–8, [OH–] = 10–6)

Note: log (10x) = x


Acid/Base Strengths

pH –log[H+]

pKa –logKa

A pH meter measures the [H+] in solution

For an acid in equilibrium with its ions:

HA(aq) H+ + A–

Equilibrium Constant Ka [H+][A–]

[H+] Molar concentration of H+

= # of moles of H+ / L of solution

[HA]

Definitions:

pH = log([A–]/[HA]) + pKa

log([A–]/[HA])

pH

pKa

y = a· x + b

extrapolation

Interactive pH measurements: http://www.chem.iastate.edu/group/Greenbowe/sections/proj

ectfolder/flashfiles/acidbasepH/ph_meter.html


Oxides as Acids & Bases

In general:

Nonmetallic oxides are acidic

Metallic oxides are basic

CO2(g) + H2O(l) ↔ “H2CO3(aq(”

“H2CO3(aq(” ↔ H+ + HCO3–

2 SO2(g) + O2(g) 2 SO3(g)

SO3(g) + H2O(l) “H2SO4(aq(”

“H2SO4(aq(” H+ + HSO4–

CaO(s) + H2O(l) Ca(OH)2(aq)

Ca(OH)2(aq) Ca2+ + 2 OH–

Nonmetal Oxides:

Metal Oxides:

“Acid Rain”: From NOx & SOx gases


Precipitation

Acid-Base

Making Sense of the Vast Variety of Chemical Reactions &

Balancing Chemical Reactions

Decomposition

or

Dissociation

(Gas-Forming)

HgO(s) Hg(l) + O2(g)

HCl(aq) + NaOH(aq)

Pb(NO3)2(aq) + K2CrO4(aq)

2

2HgO(s) 2Hg(l) + O2(g)

½

acid + base salt + water

PbCrO4(s) + KNO3(aq)

NaCl(aq) + H2O(l)


P4(s) + 6Cl2(g) 4PCl3(l) H2(g) + ½O2(g) H2O(l)

Note: fire or flames is heat + light

Mg(s( + ½O2(g) MgO(s) P4(s) + 5O2(g) P4O10(s)

Combustion, Redox, Formation

+ Oxide-Forming:


Formation Rxns. (“reactions”): Production of 1 mole of compound from its stable elements.

Al(s) + Br2 (l) Al2Br6(s)

2Al(s) + 6Br(g) Al2Br6(s)

Examples:

Formation of Al2Br6(s):

Formation of ZnS(s):

Formation of NaCl(s):

Zn(s) + 1/8 S8(s) ZnS(s)

Na(s) + Cl2 (g) NaCl(s)

2 3

½


Balancing Organic Reactions

Propane gas

C3H8(g) + O2(g) CO2(g) + H2O(l) 3 4 5

Note: The complete combustion of a hydrocarbon produces CO2 and H2O.


EXCEPTIONS SOLUBLE

COMPOUNDS

None Salts of alkali metals and of NH4

+

None

Salts of: NO3–

ClO3–

ClO4–

Ac–

Ag+, Hg22+, Pb2+ Salts of Cl–, Br–, I–

Examples:

(NH4)2S(aq) 2NH4+ + S2–

Hg2(NO3)2(aq) Hg22+ + 2NO3

–

Hg(NO3)2(aq) Hg2+ + 2NO3–

Ca3(PO4)2(s) (no change)

EXCEPTIONS POORLY

SOLUBLE COMPOUNDS

alkali metals

and NH4+

(of course)

Salts of: CO32–

C2O42–

PO43–

CrO42–

S2–

Hydroxides OH–

Oxides O2–

Solubility of Ionic Compounds General Solubility Guidelines for Some Ionic Compounds in Water


Pb(NO3)2(aq) + K2CrO4(aq) PbCrO4(s) + KNO3(aq)

Double-Exchange, Double-Replacement:

Net Ionic Reactions

Gross Rxn.:

Ionic Rxn.: Pb2+ + 2NO3– + 2K+ + CrO4

2– PbCrO4(s) + 2K+ + 2NO3–

Net Ionic Rxn.: Pb2+ + CrO42– PbCrO4(s)

Single-Exchange, Single-Replacement:

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

2H+ + 2Cl– Mg2+ + 2Cl–

Mg(s) + 2H+ Mg2+ + H2(g)

Gross Rxn.:

Net Ionic Rxn.:

Precipitation Rxn.

Gas-Forming +

Redox

Spectator Ions

2

Not all salts are relatively soluble in water. But all soluble salts dissociate into ions.


2NaNO3(aq) + K2CrO4(aq) Na2CrO4(aq) + 2KNO3(aq)

Double-Exchange, Double-Replacement ???

Gross Rxn.???:

2Na+ + 2NO3– + 2K+ + CrO4

2– 2Na+ + CrO42– + 2K+ + 2NO3

–

Net Ionic Rxn.: NOTHING: NO REACTION!!!


Acid-Base Rxns. (Double-Exchange)

Strong Acid + Strong Base:

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) Gross Rxn.:

H+ + Cl– + Na+ + OH– Na+ + Cl– + H2O(l)

H+ + OH– H2O(l) Net Ionic Rxn.:

Weak Acid + Strong Base:

HAc(aq) + NaOH(aq) NaAc(aq) + H2O(l)

Salt Water

HAc(aq) + Na+ + OH– Na+ + Ac– + H2O(l)

Net Ionic Rxn.: HAc(aq) + OH– Ac– + H2O(l)

Gross Rxn.: Salt Water


Strong Acid + Weak Base:

2HCl(aq) + CaCO3(s) CaCl2(aq) +

“H2CO3(aq)”

Gross Rxn.:

2H+ + 2Cl– + CaCO3(s) Ca2+ + 2Cl– + H2O(l) + CO2(g)

2H+ + CaCO3(s) Ca2+ + H2O(l) + CO2(g) Net Ionic Rxn.:

Strong Acid + Weak Base:

HCl(aq) + NH3(aq) NH4+ + Cl– or NH4Cl(aq)

H+ + Cl- + NH3(aq) NH4+ + Cl-

Net Ionic Rxn.: H+ + NH3(aq) NH4+

Gross Rxn.: Salt

(Double-Exchange)

Base (salt)

(Single-Exchange)

Salt Acid


Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

+[O], oxidation

-[O], reduction

reducing agent

oxidizing agent

Redox (Reduction-Oxidation) Rxns.

In the beginning, Redox rxns. were defined in terms of actual O’s or H’s transferred.

C2H4(g) + H2(g) C2H6(g)

-[H], oxidation

+[H], reduction

reducing agent

oxidizing agent

O-transfer: From Fe2O3 to CO

H-transfer: From H2 to C2H4

[O]

[H]

But later, Redox rxns. were more generally defined in terms of electron-transfer even

if O’s or H’s were not actually involved. So, check the rxns. above.

How many e’s are

transferred in

each rxn.?


2Fe(s) + O2(g) + 2H2O(l) 2Fe(OH)2(aq)

4Fe(OH)2(aq) + O2(g) 2H2O + 2Fe2O3·H2O(s)

Corrosion

Brown-red (“rust”)

(+2) (0)

(+3) (+2)

oxidized

oxidized

reduced

reduced


2Ag+ + Cu(s) 2Ag(s) + Cu 2+

–2e–, oxidation

+2e–, reduction

reducing agent

oxidizing agent

Redox (Reduction-Oxidation) Rxns. (continued)

So, Redox rxns. can be more generally defined in terms of electron-transfer

even if O’s or H’s were not actually involved.

e-transfer: From Cu to Ag+

Write 2 Half-Rxns.:

Red.: 2[Ag+ + e– Ag(s)]

Ox.: Cu(s) Cu2+ + 2e–

Total: 2Ag+ + Cu(s) 2Ag(s) + Cu2+

Oxidation: Oxidation # (“charge”( increases: e’s are lost

Reduction: Oxidation # (“charge”( decreases: e’s are gained

2e–


Koren’s “Law of the Jungle”

In the Jungle,

The Mighty Jungle,

The Lion

Sleeps Tonight …

(But when awakened he’ll roar)

LEO

GER!

LEO: Loss of Electrons is Oxidation

GER: Gain of Electrons is Reduction


Electronegativity

The ability of a bonded atom to draw electrons close to it.

Oxidation Numbers (or Oxidation States)

7A

H : H

F : F

H : F δ+ δ-

O.N.:

• is an “invention” to explain Redox rxns.

• is a make-believe, virtual reality, electron-book-keeping charge on a bonded atom.

• indicates how electrons are shared among the various atoms bonded together.

• assumes that the more electronegative atom completely steals the electron(s) in a

bond, i.e., exaggerates the bond as ionic.

Na + F Na F + – Recall:

So, “oxidation number” of F is –1 and of H is +1:

Complete e-transfer

Full Charges

Partial e-transfer Partial Charges

middle

Redox (Reduction-Oxidation) Reactions

(continued)


Rules for Determining Oxidation Numbers Note: The O.N. is the value on one atom of an element.

1. The O.N. of an atom in the molecule of a pure element is zero (0).

Examples: O in O2 or O3; S in S8; P in P4; Na in Na(s)

2. The O.N. of an atom in a monatomic ion is the same as its charge.

Examples: Cl–; Na+; Al3+

3. Some elements have the same oxidation numbers in ALL their compounds.

(a) Group IA Metals have an O.N. = +1 in all their compounds.

(b) Group IIA Metals have an O.N. = +2 in all their compounds.

(c) F has an O.N. = –1 in all its compounds.

4. Some elements have the same oxidation numbers in nearly all their compounds.

(a) H in covalent compounds is always +1; Examples: H2O; HCl; ...

H bonded to a IA or IIA metal is –1, of course. Examples: NaH; MgH2.

(b) O has an O.N. = –2 in most compounds. Examples: CO2; MgO; CH3OH ...

O as a peroxide is –1. Examples: H2O2; Na2O2,

O as a superoxide is –½. Example: KO2.

O with F has an O.N. of +2. Example: OF2.

5. The sum of all the O.N.’s equals the charge on the molecular species.

Work out the O.N. for each atom in the following species: CaCO3, SO42–, NH4

+


Summary of Types of Reactions

(alphabetical)

Acid-Base

Combustion

Decomposition, Dissociation

Double-Exchange

Exothermic

Formation

Gas-Forming

Hydrolysis

Net Ionic

Oxide-Forming

Precipitation

Redox (Reduction-Oxidation)

Single Exchange


5.

Measurement Units & Conversions


The S.I. System of Measurement Units

“Système International”, SI (1960)

A systematic modernized metric system based on 7 base units

Built from the old “MKS” system

NIST: http://physics.nist.gov/cuu/Units/units.html

http://physics.nist.gov/cuu/Units/units.html

http://physics.nist.gov/cuu/Units/units.html


Base quantity SI base unit

Name Symbol

length meter m

mass kilogram kg

time second s

electric current ampere A

temperature kelvin K

amount of substance mole mol

luminous intensity candela cd


Examples of SI derived units

SI derived unit

Derived quantity Name Symbol

area square meter m2

volume cubic meter m3

speed, velocity meter per second m/s

acceleration meter per second squared m/s2

wave number reciprocal meter m-1

mass density kilogram per cubic meter kg/m3

specific volume cubic meter per kilogram m3/kg

current density ampere per square meter A/m2

magnetic field strength ampere per meter A/m

amount-of-substance

concentration mole per cubic meter mol/m3

luminance candela per square meter cd/m2

mass fraction

kilogram per kilogram,

which may be represented

by the number 1

kg/kg = 1


SI derived units with special names and symbols

Derived quantity

SI derived unit

Name Symbol

Expression in terms of

other SI units

Expression in terms of

SI base units

frequency hertz Hz - s-1

force newton N - m·kg·s-2

pressure, stress pascal Pa N/m2 m-1·kg·s-2

energy, work, quantity of heat joule J N·m m2·kg·s-2

power, radiant flux watt W J/s m2·kg·s-3

electric charge, quantity of electricity coulomb C - s·A

electric potential difference,electromotive force (emf) volt V W/A m2·kg·s-3·A-1

capacitance farad F C/V m-2·kg-1·s4·A2

electric resistance ohm Ω V/A m2·kg·s-3·A-2

electric conductance siemens S A/V m-2·kg-1·s3·A2

magnetic flux weber Wb V·s m2·kg·s-2·A-1

magnetic flux density tesla T Wb/m2 kg·s-2·A-1

inductance henry H Wb/A m2·kg·s-2·A-2

Celsius temperature degree Celsius °C - K

luminous flux lumen lm cd·sr (c) m2·m-2·cd = cd

illuminance lux lx lm/m2 m2·m-4·cd = m-2·cd

activity (of a radionuclide) becquerel Bq - s-1

absorbed dose, specific energy (imparted), kerma gray Gy J/kg m2·s-2

dose equivalent sievert Sv J/kg m2·s-2

catalytic activity katal kat s-1·mol


Factor Name Symbol Factor Name Symbol

1024 yotta Y 10-1 deci d

1021 zetta Z 10-2 centi c

1018 exa E 10-3 milli m

1015 peta P 10-6 micro µ

1012 tera T 10-9 nano n

109 giga G 10-12 pico p

106 mega M 10-15 femto f

103 kilo k 10-18 atto a

102 hecto h 10-21 zepto z

101 deca,

deka da 10-24 yocto y

SI prefixes


Units outside the SI that are accepted for use with the SI

Name Symbol Value in SI units

minute (time) min 1 min = 60 s

hour h 1 h = 60 min = 3600 s

day d 1 d = 24 h = 86 400 s

liter L 1 L = 1 dm3 = 10-3 m3

metric ton (“tonne”) t 1 t = 103 kg

neper Np 1 Np = 1

bel B 1 B = (1/2) ln 10 Np (c)

electronvolt eV 1 eV = 1.602 18 x 10-19 J,

approximately

unified atomic mass unit u 1 u = 1.660 54 x 10-27 kg,

approximately


Other units outside the SI that are currently accepted for use with the SI,

subject to further review

Name Symbol Value in SI units

nautical mile 1 nautical mile = 1852 m

knot 1 nautical mile per hour =

(1852/3600) m/s

are a 1 a = 1 dam2 = 102 m2

hectare ha 1 ha = 1 hm2 = 104 m2

bar bar 1 bar = 0.1 MPa = 100 kPa

= 1000 hPa = 105 Pa

ångström Å 1 Å = 0.1 nm = 10-10 m

barn b 1 b = 100 fm2 = 10-28 m2

curie Ci 1 Ci = 3.7 x 1010 Bq

roentgen R 1 R = 2.58 x 10-4 C/kg

rad rad 1 rad = 1 cGy = 10-2 Gy

rem rem 1 rem = 1 cSv = 10-2 Sv


Common Units of Measurements in Chemistry

Conversion to other units Symbol Unit name Quantity

1 in ≡ 2.54 cm m

in

meter

inch Length, l

1 lb 453.6 g kg

lb

kilogram

pound Mass, m

s second Time, t

T(K) = t(oC) + 273.15 oC

K

degree Celsius,

kelvin

Temperature,

T

1 bar = 0.1 MPa = 100 kPa = 105 Pa

1 atm ≡ 1.01325 bar

1 atm ≡ 760 mm Hg

1 torr ≡ 1 mm Hg

Pa

bar

atm

mm Hg, torr

pascal

atmosphere

bar

mm Hg, torr

Pressure, P

1 mol of items 6.02 x 1023 items

(Avogadro’s number = 6.02 x 1023 ) mol mole

Amount of

substance, n

1 L = 103 mL = 10–3 m3

1 cm3 ≡ 1 mL

m3

L

mL

cm3

cubic meter

liter

milliliter

cubic centimeter (c.c.)

Volume, V

d = m/V. 1 kg/m3 = 10–3 g/cm3

[@ room temp, d (g/mL): H2O: 1.0; Hg: 13.6]

kg/m3

g/cm3 = g/mL kilogram per cubic meter Density, d or ρ

1 cal ≡ 4.18 J J

cal

joule

calorie

Energy, heat,

work


Conversions: Unit Conversion Factor

The Principle:

Multiplying a number (or measurement) by “1”

does not change the value of the number (or measurement)!!!

Examples:

5 1 = 5

(5 in) 1 = 5 in

5 𝑖𝑛 × 2

2= 5 𝑖𝑛

And now:

Convert “5.0 in” to “cm”; that is, how many centimeters are equal to 5.0 in?

The Unit-Factor for “cm in” is: 2.54 𝑐𝑚

1 𝑖𝑛 or

1 𝑖𝑛

2.54 𝑐𝑚

5.0 in 2.54 𝑐𝑚

1 𝑖𝑛 = 12.7 cm 13 cm. Why? “Significant Figures”! (See next topic)

Convert “9.0 cm” to “in”; that is, how many inches are equal to 9.0 cm?

9.0 cm 1 𝑖𝑛

2.54𝑐𝑚 = 3.5433 in 3.5 in. Why? “Significant Figures”! (See next topic)


Significant Figures

The Principle:

The result of a mathematical operation involving measurements

cannot be more accurate than the least accurate measurement)!!!

Consider the following measurements:

The distance (or length) between two specific points were measured with

different measuring devices/instruments (e.g, ruler, vernier caliper, laser, etc.):

9 cm

9.0 cm

9.00 cm

9.000 cm

Which measurement is the most accurate (least uncertainty)?

What is the number of “significant figures of each measurement?

Thus, consider the previous exercise of converting: 9.0 cm in:

9.0 cm 1 𝑖𝑛

2.54𝑐𝑚 = 3.5433 in 3.5 in. Why?

Sometimes we can write the final result as such: 3.5433 or 3.5433 or 3.5433


More Unit Conversions

Area:

This property consists of the square of a unit:

Density:

This property consists of a ratio of units:

Multiple Factors:

This property consists of a product of units:

Convert 5.00 m2 to cm2:

5.00 m2 100 𝑐𝑚

1 𝑚

2 = 5.00 100 𝑐𝑚 2 = 5.00 (102 cm)2 = 5.00 104 cm2

Volume :

This property consists of the cube of a unit: Convert 5.00 m3 to cm3:

5.00 m3 100 𝑐𝑚

1 𝑚

3 = 5.00 100 𝑐𝑚 3 = 5.00 (102 cm)3 = 5.00 106 cm3

Convert 5.00 m to ft: 100 𝑐𝑚

1 𝑚

1 𝑖𝑛

2.54 𝑐𝑚

1 𝑓𝑡

12 𝑖𝑛 = 16.4 ft 5.00 m

Convert density of Hg, 13.6 g/cm3, to kg/m3:

13.6 = 13.6 103 kg/m3 𝑔

𝑐𝑚3

1 𝑘𝑔

103 𝑔

102𝑐𝑚

1 𝑚

3


6.

The Mole


mole

Mole Day is celebrated on Oct. 23 from 6:02 am to 6:02 pm


Lorenzo Romano Amadeo Carlo Avogadro 1776 – 1856

The word “mole”:

Coined by Wilhelm Ostwald (1896), Latin “moles” meaning “heap” or “pile”.

Avogadro’s Number:

1 mole of things = 6.022x1023 things


The mole is a number!

For example:

1 dozen = 12

1 dozen eggs = 12 eggs

1 dozen chairs = 12 chairs

1 dozen molecules = 12 molecules

1 dozen atoms = 12 atoms

1 mole = 6.022x1023

1 mole of eggs = 6.022x1023 eggs

1 mole of chairs = 6.022x1023 chairs

1 mole of molecules = 6.022x1023 molecules

1 mole of atoms = 6.022x1023 atoms

The Mole … once agin … The Bridge to the Human Scale


Molecular Formula, Moles, Molecules, Atoms: Summary

Examples:

Molecular formula = (CH3)2CF2

1 molecule contains: 3 C atoms

6 H atoms

2 F atoms

100 molecules contain: 300 C atoms

600 H atoms

200 F atoms

1 mole of molecules contains: 3 moles of C atoms = 3 x 6.022 x 1023 C atoms

6 moles of H atoms = 6 x 6.022 x 1023 H atoms

2 moles of F atoms = 2 x 6.022 x 1023 F atoms

OK? OK!


What’s so great about the mole or Avogadro’s Number???

Converts masses in amu to the same number in grams

Molar Mass

(g/mol) Masses Atoms/Molecules

12 g/mol

12 amu

12 g

1 12C atom

1 mol of 12C atoms

10.0129 g/mol

10.0129 amu

10.0129 g

1 10B atom

1 mol of 10B atoms

32.00 g/mol

32.00 amu

32.00 g

1 O2 molecule

1 mol of O2 molecules

2.00 g/mol

2.00 amu

2.00 g

1 H2 molecule

1 mol of H2 molecules

Molecular Weight (MW)

Formula Weight (FW)

1 g = 6.022x1023 amu

g ? amu

Avogadro’s Number is a Magic Number!!!

Atomic Weight (AW)


1 g = 6.022x1023 amu

g amu

5.0 g amu = ?

5.0 g g amu

1 = 3.0 x 1024 amu

“Unit Factor”

Mass Conversions

6.022 x 1023

Example for g amu:

2.0 amu g = ?

2.0 amu g

amu 1 = 3.3 x 10-24 g

“Unit Factor”

6.022 x 1023

Example for amu g:


Grams Moles 31.8 g Cu Moles of Cu atoms = ?

31.8 g Cu g Cu

mol Cu 1

63.55 = 0.500 mol Cu

Moles Grams

“Unit Factor”

(A.W.)

0.300 mol Cu atoms g Cu = ?

g Cu

mol Cu 1

63.55 = 19.1 g Cu

“Unit Factor”

(A.W.)

0.300 mol Cu

Moles and Molar Masses of the Elements


How many moles are there in 1.00 lb of silicon?

Weight

1.00 lb Si g

lb 1

453.6 mol g 28.0855

1 = 16.2 mol Si

Moles

Grams Moles Weight


Grams Moles 1.6 g CH4 Moles of CH4 molecules = ?

1.6 g CH4 g CH4

mol CH4 1

16.0 = 0.10 mol CH4

Moles Grams

“Unit Factor” (M.W.)

0.10 mol CH4 molecules g CH4 = ?

g CH4

mol CH4 1

16.0 = 1.6 g CH4

“Unit Factor” (M.W.)

0.10 mol CH4

Molar Mass or Molecular Weight (M.W.) For Example, for CH4: MW = 12.0 + 4(1.0) = 16.0 g/mol

Also:

M

mn


Moles Number

Number of Items

0.20 mol tables # of tables = ?

0.20 mol tables mol tables

tables 6.02x1023

1 = 1.2x1023 tables

Number Moles

“Unit Factor”

1.2x1023 Cu atoms mol Cu = ?

mol Cu atoms

Cu atoms 6.02x1023

1 = 0.20 mol Cu atoms

“Unit Factor”

1.2x1023 Cu atoms

Avogadro’s Number:

1 mole of things = 6.022x1023 things


Percent Composition

Weight (or mass) % of each element in a compound

For example, NH3:

N % 82.27 100 x g 17.030

g 01.14 NHin N % 3

H % 17.73 100 x g 17.030

g 1.008 x 3 NHin H % 3

100.00 %


Summary

Mole Avogadro’s

Number Mass

constant variable


7.

Stoichiometry: Chemical Mathematics


Stoichiometry Stoicheion + metron

(element) (measure)

Weight relations in chemical rxns. based on conservation of matter

Examples:

2 “molecules” 1 molecule 2 atoms

2x 6.02x1023 “molec.” 6.02x1023 molecules 2x 6.02x1023 atoms

2 moles of “molecules” 1 mole of molecules 2 moles of atoms

80.6 g 32.0 g 48.6 g

2Mg(s) + O2(g) 2MgO(s)

For any rxn.,

The absolute value of each coefficient is meaningless by itself!

BUT, the RATIOS are HOLY!!!

IFThen: אז-אם relationship

כימות כימי


Stoichiometric Calculations: The Approach

gram mole mole mole

moleA gram gram gramA

MW (or AW)

2. Think in Moles !!!

3. Setup a Flow-Chart whereby g mol mol

Helpful Tips for Solving Problems:

4. Always include Units and Substance Name

Simple formula: MW

mn

moleB Stoichiometry

(rxn)

1. Write the Balanced rxn!!


2Mg(s) + O2(g) 2MgO(s)

Stoichiometric Calculations: Examples

Calculate the number of grams of MgO produced from 0.145 g Mg.

grams

moles moles

grams ?

AW, MW MW

Stoichiometric Ratio: Rxn

0.145 g Mg

Mg 3050.24

Mg 1

g

mol

Mg 2

MgO 2

mol

mol

MgO 1

MgO 0.30444

mol

g = 0.240 g MgO

MW factor MW factor Rxn factor or

Stoichiometric factor

So, Remember, All Roads Go Through Moles !!!


Percent Yield

% Yield = 100 x

yieldltheoretica

yieldactual

Many rxns do NOT go to completion. There is a chemical energy barrier involved.

C7H6O3(s) + C4H6O3(l) C9H8O4(s) + H2O(l) For example:

salicylic acid

acetic anhydride

acetyl salicylic

acid

Aspirin

If from 14.43 g of the acid, 6.26 g of aspirin is produced, what is the % yield for the rxn?

2 2

Theoretical yield or maximum yield (assuming rxn goes to completion):

cid 38.12261

cid 1

ag

amol

acid 2

aspirin 2

mol

mol

aspirin 1

aspirin 80.15981

mol

g = 18.82 g aspirin

% yield = 100 x 82.18

26.6

g

g= 33.3 %

14.43 g acid


Solution Concentrations

Molarity, Formality, molality, Normality, % w/w, % w/v

Molarity = M = n

/V # of moles of solute per liter of solution

# of mmoles of solute per mL of solution

moles/L

mmoles/mL

“3.0 molar potassium permanganate solution”

3.0 moles of KMnO4 per L of solution

3.0 mmoles of KMnO4 per mL of solution

Dilutions

Prepare 500.0 mL of a 0.100-M KMnO4 solution from a 3.0-M stock solution?

Use n = M•V

= n = Mstock • Vstock Mnew • Vnew

Vstock = 0.0167 L = 16.7 mL

Preparation of a Diluted Solution:

• Remove 16.7 mL of the 3.0-M stock solution

• Add enough water (“483.3 mL”) to produce 500. mL of solution


Solution Stoichiometry

Titrations: Acid-Base & Redox

Buret

NaOH(aq) + H2C2O4(aq)

Acid-Base Titrations

Na2C2O4(aq) + H2O(l) 2 2

Erlenmeyer

flask

End-Point or Equivalence Point of an Acid-Base or Redox rxn: Point at which all of of the Acid reacts with all of the Base or all of the oxidant reacts with all of the reductant.

Problem 1: Reaching an end-point

How many mL of 0.300 M sodium hydroxide are needed to

titrate 25.0 mL of 0.400 M oxalic acid?

Solution:

n = M ·V

MA,VA moles acid moles base V base rxn V=n/M

nACID = MA ·VA =

VBASE = nBASE / MBASE = 0.0200 mol / 0.300 M = 0.0667 L = 66.7 mL

(0.400 M)(0.0250 L) = 0.0100 mol

nBASE = 2 • nACID = 2 • (0.0100 mol) = 0.0200 mol

n=M·V


Terminology for dissolution: a solute is dissolved by a solvent to form a solution.

Units Formula Definition Name

M M = molssolute/Lsol’n # of moles of solute per L of solution Molarity

N N = eqssolute/Lsol’n # of equivalents of solute p. L of sol’n Normality

g/mL dsol’n = msol’n/Vsol’n mass of the sol’n per volume of sol’n Density of Sol’n (d or , rho)

m m = molssolute/kgsolvent

|m| |M|

# of moles of solute per kg of solvent

In dilute aqueous solutions: Molality

(none)

%

Xi = ni/nsol’n, ΣXi = 1

Xi 100

moles of a component p. total moles

mole fraction of a comp. as a percent

Mole fraction

Mole percent

% w/w % = mi/msol’n100

w/v % = gi/mLsol’n100

weight of solute/weight of sol’n, as %

weight of solute/volume of sol’n, as % Weight percent

ppm ppm = gi/gsol’n= mg/kg

ppm |mg/L|

# of solute parts p. million sol’n parts

In dilute aqueous solutions: Parts per million

ppb ppb = ngi/gsol’n= g/kg

ppb |g/L|

# of solute parts p. billion sol’n parts

In dilute aqueous solutions: Parts per billion

% v/v %=Vi,pure/Vsol’n100 vol. of pure solute/vol. of sol’n, as % Volume percent

Proof Proof = 2(v/v %) Double the Volume % (for whiskey) Proof

More Solution Concentrations

Note: Must always write the units and the substance, e.g., 2.0 g solute.

Questions: What is ppt? ___________________ pph? ________________________


Selected Concentration Examples:

(1) 1.2 kg ethylene glycol (HOCH2CH2OH), an antifreeze, is added to 4.0 kg

water. Calculate (for ethylene glycol): mole fraction, molality, weight/weight%.

[Answers: X = 0.080, m = 4.8 m, w/w % = 23 %]

(2) 560 g NaHSO4 are dissolved in 4.5x105 L water at 25 oC. Calculate the Na+

concentration in parts per million. [Answer: 0.24 ppm]

(3) 10.0 g of sucrose (C12H22O11) are dissolved in 250. g of water. Calculate (for

sugar): X, m, w/w %. [Answers: X = 0.00210, m = 0.117 m, w/w % = 3.85 %.]

(4) Sea water has a sodium ion concentration of 1.08 x 104 ppm. If the Na is

present in the form of dissolved sodium chloride, how many grams of NaCl

are in each liter of sea water? (Density of sea water is 1.05 g/mL.)

[Answer: 28.7 g NaCl/L]

(5) A 0.100-M aqueous solution of ethylene glycol has a density of 1.09 g/mL.

What is the molality of the solution? [Answer: 0.0923 m]


Impossible (?) Formulas

HIJKLMNO

BaNa2

NAg

MgNiV

13Al–1H–19K–9F–19K

S H N U K Ra

Au

:(L R)ארגון אשלגן חנקן הליום גופרית

c- he-m- i= s-tr y for conservators · lead pb ancient latin, plumbum, meaning heavy french,...

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